laboratory manual for inorganic chemistry …€¦ · · 2017-11-23the experiment. no exceptions...

LABORATORY MANUAL

FOR

INORGANIC CHEMISTRY

FIRST EDITION

DEPARTMENT OF CHEMISTRY

BABCOCK UNIVERSITY. ILISAN. REMO. OGUN STATE

2014

CONTENTS

Policy, Practices, and Procedures iii

EXPERIMENTS

1. Precipitation Reactions

2. Redox and Other Chemical Reactions

3. Equilibria and Le Charteliers Principle

4. Estimating Equilibrium Constant

5. Gravimetry: Determination of Oxalates

6. Limiting Reactants

7. Complex Reactions

Policies, Practices, and Procedures

SAFETY

The most important events in the laboratory are safety and the right to know the physical and

chemical hazards present in the laboratory. Follow instructions as stated in the laboratory manual

and be careful with acids and bases.

!!!PLEASE DO NOT DISOBEY LABORATORY RULES. IT MIGHT PROVE FATAL!!!!

PROTECTION

1. Instructors, laboratory technologist, laboratory assistants, students and other personnel

that work in the laboratory must wear approved goggles at all times when in the

laboratory. It is a recommendation not to wear contact lenses when in working in the

laboratory

2. It is advisable to wear old clothes that cover heat to toe when working in the laboratory.

Long sleeve shirts and long pants are recommended. Shorts, midriff tops and sandals are

not allowed.

3. Long hair is to be tied and well tucked away. It is a fire hazard when let loose or flying

4. Students are not allowed to work in the laboratory alone

5. Horseplay or other pranks are prohibited in the laboratory

6. Smoking, eating, drinking or applying makeup is not allowed in the laboratory

LABORATORY MAINTENANCE

1. Make sure your Laboratory space(s) is cleaned. Also clean all the equipments and

returned them to their assigned positions. Failure to do so will result to a zero grade for

the experiment. NO exceptions please.

2. All glassware must be cleaned before it is put away.

3. Use sponges to clean bench tops and wiping of non-hazardous materials.

4. Laboratory instructors are the ONLY one allowed to clean up corrosive or toxic materials

5. Sweep up broken glassware with a broom and collect with the dust pan and then place in

the special container provided for glasses.

6. No debris of any type should be left in the sink. Put all debris in allocated containers

7. Make sure all drawers are properly closed and locked when necessary.

GENERAL INFORMATION

1. Dispense organic solvents, strong acids and bases and other volatile solvents in the fume

hoods.

2. No fee will be collected for broken equipment or glassware. Each broken glassware

and equipments will be replaced with two of similar type by the culprit.

LABORATORY TECHNIQUES

1. Use proper utensils such as crucible tongs to hold or move hot items.

2. Make sure there are no flammable materials near you when lightning a burner

3. Add boiling chips to liquids before heating them up. This will help to prevent bumping or

boil over.

4. Place test tubes in a slanting position away from yourself and others when heating

liquids. Heat liquids at the surface of the liquid.

5. Do not heat up a closed system

6. Heat all substances that emit noxious fumes under the hood

7. Use funnel to transfer liquids into a narrow neck container

8. Use a bulb or pump to pipette a liquid. Never use your mouth

9. Avoid smelling anything unless instructed to do so. While sniffing, gently waffle the

material towards your nose when allowed to do so

10. Do not return excess reagent to its original container

11. Do not experiment with the chemicals in the laboratory except those that your are

scheduled to do.

12. Do not use your pipette or spatula to remove samples from the stock container. Use the

one provided by the laboratory technologist

13. Correctly label test tubes or other containers indicating their contents

14. Strong acids and bases should be added to water and not vice versa.

EMERGENCIES AND FIRES

1. Laboratory instructors are in-charge of all emergencies. Follow instructions as directed

2. All laboratory users should learn how to locate the following materials: safety shower,

eyewash, blankets, fire extinguishers, first aid kit, fire alarm

3. Laboratory users should notify the laboratory instructors of any fire.

4. Turn off all gas jets if it is the source of the fire

5. All laboratory users should learn how to use the fire extinguisher

ACCIDENTS AND INJURIES

1. The chemistry department does not treat injuries or illness. Any injury or illness will be

referred to the University Medical Center, Ilishan, Ogun State.

2. It is the responsibility of the laboratory instructor(s) on duty to prevent further injury by

taking the appropriate action after the incident. Arrangement should be made to

immediately transport the victim to the Medical Center. If the injury is minor and the

student can walk to the Medical Center, such student should be accompanied by another

person to the Medical Center.

3. An accident report form must be filled at all times even when the victim declines Medical

treatment.

SPECIAL WASTE

1. The laboratory will provide label containers for hazardous waste. Read the label very

well and dispose the waste appropriately.

2. At no time should organic or toxic wastes such as mercury, lead, chromium be dumped

down the drain.

3. Ask when in doubt about proper disposal of waste.

Experiment 1: Precipitation Reactions

Theory

Many methods are used in determining the end points in precipitation reactions. For example, Mohr

titration reactions are employed in the determination of chlorides and bromides. In the titration, a neutral

chloride solution is titrated with a silver nitrate using potassium chromate as an indicator. At the end

point, the chromate ions combine with the silver ions for form sparingly soluble red silver chromate. In

this experiment, we have two sparingly soluble salts being silver chloride (Ksol= 1.2 x 10-102) and silver

chromate (Ksol= 1.7 x 10-12) respectively. If for example,0.1 M sodium chloride in the presence of dilute

potassium chromate is titrated with 0.1 M silver nitrate, silver chloride is a less soluble salt and again the

chloride concentration is high, hence silver chloride will be precipitated. At the first appearance of the red

silver nitrate, we shall have both salts in equilibrium in solution hence;

[Ag+][Cl-] = KsolAgCl= 1.2 x10-10 and [Ag+][CrO4-] = KsolAgCrO4= 1.7 x 10-12

[Ag+] = KsolAgCl/[Cl-]= {KsolAgCrO4/[CrO4-]}1/2

[Cl-]/ CrO4-]= KsolAgC/ KsolAgCrO4= 1.2 x10-10/1.7 x-12= 9.2 x 10-5

At equivalence point [Cl-] ={ KsolAgCl}1/2= 1.1 x 10-5

[CrO4-] = {[Cl-]/9.2x 10-5} = {1.1 x 1—5/9.2 x 10-5} = 1.4 x 10-5

If a more dilute potassium chromate solution is used as an indicator, it makes detection of the silver

chromate more difficult (end point). This will introduce an error which can be corrected by doing a blank

determination. The blank will help us to determine the volume of silver nitrate required to give a

perceptible coloration when added to distilled water containing the same amount on indicator. This

volume is subtracted from the volume of the standard solution used.

Apparatus: Weighing balance, filter paper, 2x 100 and 2x 500 ml graduated flask, stirrer, 3x250mL

flask, 25mL pipette, 1.0 mL pipette, indicator dropper

Reagents: Pure Silver nitrate, Sodium Chloride, potassium chromate, potassium dichromate, calcium

carbonate

Procedure

Section A: Preparation of 0.1M Silver Nitrate Standard Solution

Dry some finely powered silver nitrate at 120oC for 2 hours. Allow to cool down and accurately

weigh out approximately 8.5g of silver nitrate and dissolve in 100 mL of water in a 500ml

graduated flask. Make up the solution the solution to the mark and stir the solution thoroughly.

Section B: Preparation of 0.1M Sodium Chloride Standard Solution

Weigh out 2.922 g of pure and dry sodium chloride and dissolve in a 100 ml flask in a 500 mL

graduated flask. Dilute to the mark and stir the solution thoroughly for proper mixing.

Section C: Preparation of the Potassium Chromate Indicator Standard Solution

4.2 g of potassium chromate and 0.7g of potassium dichromate are dissolved with water in a

100ml graduated flask. Dilute to the mark and stir very well for proper mixing

Section D: The Mohr Titration (Titration of the Silver Nitrate Solution)

Pipette 25 mL of the standard sodium chloride solution into a 250 mL conical flask resting on a

white tile and add 1.0 mL of the indicator with the 1.0 mL pipette. Pour the silver nitrate into a

burette and gently add to the sodium chloride solution, swirling continuously until a red color

formed by the addition of each drop of the silver nitrate begins to disappear slowly. Continue the

addition drop wisely until a faint but distinct color change occurs. The faint reddish brown color

should persist after a brief shaking. This is the end point. If the end point is overstepped,

production of a deep reddish-brown color) add more chloride solution and titrate again. Subtract

the blank volume from the titration volume (this should not be more than 0.03-0.10 mL) before

calculation. Repeat the titration two more times.

Section E: The indicator Blank Correction

Add 0.5 g of calcium carbonate to water that have the same volume of the volume obtained in

the silver nitrate titration. Add 1.0 mL of the indicator to the blank solution. Gradually add

0.01M solution of silver nitrate to the solution until an inert white precipitate similar to that

obtained in the titration of chlorides

Experiment 1 RESULT

Name Department Date

Table 1

Volume 1 Volume 2 Volume 3

Ag(NO3)

[Cl-]

[CrO4]

[Ag+]

Question

1. Why must be reaction carried out in a neutral solution?

2. What reaction will happen if the reaction is carried out in an acid solution?

3. What might happen if the solution is markedly alkaline?

Experiment 2: Redox and Other Chemical Reactions

Theory

Double displacement reaction simply involves exchange of ions which can be represented in

the following equation.

AB + CD → AD + CB

Acids, bases and salts commonly undergo displacement reactions because they dissociate in

aqueous medium. The dissociation can be represented thus;

A+ + B- + C+ + D- +H2O → AD + B- + C+ + H2O

Some of the ions may appear in the same form on both sides of the equation and therefore,

omitted from the net ionic equations as shown below

A+ + D- → AD

There are evidences that show that displacement reaction has occurred. For example,

precipitates may be formed, gas may be evolved (bubbles), or heat may be absorbed or

generated. Absence of any of these evidences may indicate that the possible products are

soluble in aqueous solution; hence no net reaction has occurred.

Redox Reaction: Redox is an abbreviation for reduction and oxidation reaction. Both occurs

simultaneously and abound in nature. In redox reactions, electrons are transferred unlike

during displacement reactions where ions are transferred. In essence, what is happening is

that one species of reagent looses electron while the other gain the same electron at the same

time. So, equal moles of electrons is lost and gained at the same time. Redox reaction can

then be represented by two half equations thus

A → A+ ne- - oxidation

B + ne- → B- - reduction

Apparatus: Test tubes, droppers

Procedure: Mix five drops of each of the following reagent(s) in a test tube. Mix well and note

your observation in the Table 2.1

1. 0.25M (NH4)2 C2O4 + 0.1M CoCl2 2. 0.1M Ca (NO3)2 + 0.1M NiCl2

3. 0.1M Ca (NO3)2 + 1M Na2CO3 4. 0.1M CuSO42 + 0.1M Na3PO4

Experiment 2 Cont’d

5. 0.1M Fe(NO3)3 + 6M NaOH 6. 6M HCl + 1M Na2CO3

6M HCl + 6M NaOH 8. 0.1M NiCl2 + 1M Na2CO3

Use Table A in Appendix A as a guide in determining what type of reaction is occurring in

these reactions and the net ionic equation for these reactions (if any)

Redox Reactions

Mix the following reagents in the order shown in the equations below

1. Cu + 6M HNO3

2. Mg + 6M HCl

3. 0.1 M KMnO4 + 6M HCl

4. 0.1 M KMnO4 + 6M H2SO4 + Fe SO4

Record your observations and the net ionic equation in Table 2.2

Questions

1. After adding five drop so 6M Nitric acid are added to an equal volume of Ba(OH)2, the

resulting solution becomes warmer. Write a balanced net ionic equation to describe the

reaction.

2. A strip of aluminum metal was dropped in 6M HCl and bubble so Hydrogen gas was

observed. The resulting solution contains Aluminum ions (Al3+). Write the net ionic

equation for this reaction

Experiment 2 RESULT


Table 2.1

Reagents Evidence of Reaction Net Ionic Equation

(NH4)2 C2O4 + CoCl2

Ca (NO3)2 + NiCl2

Ca (NO3)2 + Na2CO3

CuSO42 + Na3PO4

Fe(NO3)3 + NaOH

HCl + Na2CO3

HCl + Na OH

NiCl2 + Na2CO3


Table 2.2

Reagents Evidence of Reaction Net Ionic Equation

Cu+ NO3-/ Cu2+ NO

Mg +H+ / Mg2+ + H2

MnO4- + Cl-/Mn2+ + Cl2

MnO4- + Fe2+/Mn2+ + Fe3+

Experiment 3: Equilibra and Le Charteliers Principle

Theory

In a chemical reaction, reactants are consumed to make products. In some cases, all the reactants

are consumed to make the product (reactions go to completion)

NaOH + HCl → NaCl + H2O (1)

and in some cases, not all the reactants are consumed to make the product (partial reaction).

H2O ↔ H+ + OH- (2)

Another example of a partial reaction is the production of hydrogen iodide(at room temperature)

shown in the equation below. In such reactions, the rate of the forward reaction (product

formation) and the rate of the

H2 + I2 ↔ 2HI (3)

backward reaction (reactant formation) will be the same after some time. That stage is called a

dynamic equilibrium state. If the ratio of the concentration of the product with that of the

reactant is calculated, the ratio will give a constant value at that given condition (temperature,

concentration, pressure, heat of reaction etc). That ratio is called the equilibrium constant (Keq)

Keq. = [HI]2/[H2][I2] (4)

Le Charterliers Principle

Le Charteliers principle talks about how equilibrium responds to changes (stress) that happens to

a system at equilibrium. He stated that the effect of any changes made to a system at equilibrium

is annulled (cancelled) by the equilibrium re-establishing itself or the equilibrium position will

shift so as to minimize the effect of such changes. For example, in forming HI (eqn. 3), if the

concentration of H2 is increased after equilibrium is established, equilibrium will respond by

consuming more of I2 in the same proportion to produce more of HI, hence equilibrium will

be re-established.

Another example is nitrogen dioxide gas which dimerizes to form dinitorgen dioxide in an

exothermic reaction process

NO2 (g) + NO2(g) ↔ N2O4(g)

If extra heat is added to the system, the backward reaction, which is endothermic, will be favored

and equilibrium will shift towards the left (reactants, nitrogen dioxide) to re-establish

equilibrium. The converse is also true for this reaction


Apparatus: 3 test tubes, plastic droppers, brush

Reagents: Water, 0.1 M NiCl2, 6M NH3, 6M HCl, 1M MgCl2,1M MgCO3, 1M NH4Cl,

0.2%phenolphtalein, 6M NaOH, 0.1% bromocresol

Procedure

Part A. Complex Ion Formation with a Metal

Rinse a test tube with water. Add 1 ml of 0.1 M NiCl2 solution to the test tube. Record the color

in Table 11a. Add 2 drops of 6M NH3 solution the test tube and record your observation. Finally

add 2 drops of 6M HCl and note the color. Explain your observations with Le Charteliers

principle

Ni2+ + NH3 ↔ Ni (NH3)62+

Part B: Alteration of Equilibrium with a Common Ion

Equilibrium can be altered by adding an ion that is common to the reactant or the product. This is

called the Common ion Effect. A common ion is used to dissolve the precipitate of a sparingly

soluble salt.

(i) Put 10 drops of 1M MgCO3 in a clean test tube. Then add 5 drops of 1M MgCl2. Record

your observation in Table 11b and explain the reaction with the principle of common

ion effect and Le Charteliers principle

MgCO3↔ Mg2+ + CO32-

(ii) Add 5 drops of 1M MgCl2 to a clean test tube. Put 2 drops of 6M NH3 to the solution in

the test tube to produce a gel-like precipitate. Then add 5 drops of 1M NH4Cl and

shake the mixture very well.

MgCl2 + NH3 +2H2O ↔ Mg(OH)2 +NH4+

Observe what happens and record it in Table 11c. Explain your observation using common ion

effect and Le Charteliers principle.

Experiment 3 Cont’d RESULT


Table 3a

Reaction Observation

Color of Ni2+

Color of Complex

Color after adding HCl

(b) Explain your observations with Le Charteliers principle

Table 3b

Reaction Observation/Net Ionic Reaction

10 drops of1M MgCO3 in a test tube

10 drops of 1M MgCl2 to the test tube

(c) Explain your observations using common ion effect and Le Charteliers principle



Table 11c

Reaction Observation/Net Ionic Reaction

5 drops of1M MgCl2 in a test tube

2drops of 6M NH3 to the test tube MgCl2 + NH3 +2H2O ↔ Mg(OH)2 +NH4+

Add 5 drops of 1M NH4Cl and mix well

(d) Explain your observations using common ion effect and Le Charteliers principle

Question

1. A reaction is already at equilibrium and was subsequently heated. Can you predict how

equilibrium will respond? If not, what other information would you need to be able to predict

which direction equilibrium will shift?

Experiment 4: Estimating Equilibrium Constant

Theory

Nature is filled with many different types of reactions. Some reactions occur at very slow rate

(nature making fossil fuel), some at moderate rate (rusting of nails), and some at a very fast rate

(explosions). Many of these reactions go to completion and some do not. For the reactions that

do not go to completion (eqn. 1), equilibrium will be established at which the rate of formation

of the product (forward reaction) will be the same as the rate of re- forming the reactants (the

backward reaction). All equilibrium is established at a specific condition such as temperature and

pressure (for gaseous system).

A + B ↔ D (1)

The equilibrium constant for equation 1 is written as

Ke = [Product]/[Reactant] = [D]/[A][B]

It is easy to estimate Ke since the initial concentration of the reactants and the equilibrium

concentration of the reactants and product can be calculated. For example, if the initial

concentration of A and B is M moles and that of the product is x moles at equilibrium, then the

equilibrium concentration of A and B will be (M-x) moles. Therefore,

Ke = [x]/[M-x][M-x]

There are many ways in which we can measure the equilibrium concentration of the product.

One of such ways is using Ultraviolet- Visible spectrometry (UV-Vis) if the product is colored.

In UV-Vis, the amount of light absorbed by the product is proportional to the concentration of

the product as expressed in Beer-Lamberts law.

A= abc

A= Absorbance, a = molar absorptivity, b= cell length, c= molar concentration

To be able to determine the concentration of the product, a calibration curve must be obtained

with the UV-Vis spectrometer using known concentration of the product. A plot of the

absorbance of the product at different concentrations gives a straight line curve which can be

used to estimate the unknown concentration of the product. The slope of the curve is ab and can

be used to calculate the molar absorptivity (a) of the product.


Procedure

Part A: Calibration Curve

Make a 10 mL stock solution of Ferrous Thiocyanate ([FeSCN]2+] by adding 5.0 mL of 0.2 M

Fe(NO3)3, 1.0 mL of HNO3, and 4.0 mL of water in a 20 ml volumetric flask. Mix thoroughly

with a glass rod. The concentration of Fe3+ is much greater than that of the SCN-, hence it is

assumed that equilibrium will favor product formation and all the SCN- is consumed.

Calibration standards are prepared with the following dilution

1. 5 mL of stock solution + 5 mL of water



4. 1.0 mL of stock solution + 9 mL of water

Measure the absorbance or transmittance of the calibration standards (A= 2-log T) at 447nm.

0.002M Fe(NO3)3 can be used as the reference standard. Absorbance is plotted against the

concentration to obtain a calibration curve.

Part B: Estimating Equilibrium Constant

Prepare a stock solution for the estimation of the equilibrium constant as follow: 35 mL of

0.002M Fe(NO3)3 in 1M of HNO3 and 20 M 0.002M KSCN. Then prepare the following

concentration from this stock solution

Solution # 0.002M Fe(NO3)3

(mL)

0.002M KSCN

(mL)

H2O (mL)

5 5 1 4

6 5 2 3

7 5 3 2

8 5 4 1

9 5 5 0


Mix the mixtures thoroughly and measure its absorbance or transmittance and that of the

reference solution at 447 nm.

Waste: All waste can be flushed down the drain

Experiment 4 Cont’d: RESULT


Part A: Calibration Curve

Table 19.1 Solution #

[FeSCN2+] 1 2 3 4

% T at 447nm

Absorbance

Table 4.2

Part B: Estimating the Equilibrium Constant

Solution #

5 6 7 8 9

Initial mol

Fe(NO3)3

Initial mol

KSCN

Fe3+ + SCN- ↔ FeSCN2+

Ke = [FeSCN2+]/[ FeSCN2+ / [Fe3+] [SCN-]

Experiment 4 Cont’d: RESULT


Table 4.3

At Equilibrium

5 6 7 8 9

%T

Absorbance

[FeSCN2+]

from graph

Mol FeSCN2+

Mol SCN-

Mol Fe3+

[SCN-]

Ke

Average Ke =

Questions

1. Calculate the initial number of moles of a Fe3+ and SCN- prepared from 0.004M

Fe(NO3)3 and 0.006M of KSCN?

2. Convert 50% T to absorbance

3. Calculate the equilibrium constant (Ke) for the reaction shown below

Fe3+ + SCN- ↔ [FeSCN]2+]

If the initial [Fe3+] = 3x 10-2mmol and [SCN-] = 1 x 10-4mmol and the equilibrium

amount of [FeSCN]2+] = 2 x 10-5 mmol?

Experiment 5: Gravimetry: Determination of Oxalates

Theory

Gravimetric analysis is also known as quantitative analysis by weight is a process of isolating

and weighing of an element or a definite compound of the element in as pure form as possible.

The element or compound is separated from a weighed portion of the substance being examined.

A significant part of the process involves of the element to a pure stable compound suitable for

weighing. The weight of the element is then calculated from knowledge of the formula of the

stable compound and the atomic weights of the constituting element. The separation of the

element may be done in a number of ways that includes (i) precipitation (ii) volatilization (c)

electroanalytical (d) extraction and chromatography.

A demonstration of this method will be exemplified by the determination of oxalates as calcium

oxalate or calcium carbonate. Oxalates are important in medicine because of their involvement in

the occurrence of the disease called kidney or gall bladder stones. The neutral solution of alkali

oxalates can be acidified, heated to boiling and precipitated with boiling calcium chloride

solution. The precipitate can be allowed to stand for 12 hours, filtered, washed with hot water

dried and weighed as calcium oxalate or calcium carbonate or oxide.

Apparatus: Weighing balance, 2x 250 mL conical flask, measuring cylinder, filter paper

stirring rod or spatula, filtering crucible.

Reagents: Hot and cold water, calcium chloride, ammonium chloride, ethanol, anhydrous

diethyl ether, Bunsen burner or a heating mantle

Procedure

Weigh 2.0 g of calcium chloride and 0.5 g of ammonium chloride. Dissolve the ammonium chloride in 50

mL of cold deionized water and the calcium chloride in 50 ml of cold water. Heat the calcium chloride

solution to boiling and gradually add to the ammonium chloride solution. Allow the mixture to cool

down and treat with one third of its volume of 90 percent ethanol. Allow the mixture to stand for 30

minutes. Pour the mixture through a filtering crucible fitted with a filter paper and wash the precipitate

with warm water (50-60oC) until the chloride reaction is negative. Again wash the precipitate with cold

water(20 mL) one time, five times with ethanol and several times with small volumes (5mL) of dry

diethyl ether. The precipitate is air dried for about ten minutes and allowed to cool down in a desiccator

for another ten minutes. The precipitate is then weighed as calcium oxalate monohydrate (CaC2O4.H2O)

or oven dried until a constant weight is obtained and weighed as the carbonate or oxide.

Experiment 5: Cont’d RESULT


Table 5

Data First Attempt Second Attempt

Initial weight of Calcium

Chloride

Final weight of the Oxalate or

carbonate of oxide

% of oxalate ion ppt.

Questions:

1. Calculate the percentage of oxalate ion in 5.0 g calcium oxalate

2. What happens when an oxalate ion like calcium oxalate is heated?

3. Use an equation to explain the reaction in question 2?

Experiment 6: Limiting Reactants

Theory

The stoichiometry of any reaction determines the maximum amount of products that such

reaction will produce. For example, in the equation shown below,

Na2CO3 + Ba(C2H3O2) 2 → 2NaC2H3O2 + BaCO3

one mole of sodium carbonate reacts with one mole of barium acetate (reactants) to produce two

moles of sodium acetate and one mole of barium carbonate (products). If in the beginning, half

mole of sodium carbonate was used, then all of it will be consumed and one mole of sodium

acetate will be produced and similarly if half mole of barium acetate was present in the

beginning, then half mole of barium carbonate will be produced. In summary, we can say that the

amount of product produced will be determined by the amount of reactants in short supply at the

beginning of the reaction. The reactants in short supply at the beginning of the reaction is called

the limiting reactant.

This experiment is going to use a mixture (unknown amounts) of sodium carbonate and barium

acetate which will be dissolved in water. The two reactants will dissolve in water and produce

the product shown in the equation above. The amount of barium carbonate produced will be

determined by the amount of the limiting reactant present in the unknown mixture. By weighing

the precipitate, one will be able to determine the amount of the limiting reactants present in the

unknown mixture. One can also determine the amount of the excess reactant by adding more of

the limiting reactant followed by separation, drying and weighing.

Apparatus: Rubber tubing’s, Buchner funnel and flask, filter paper, water tap with suction

point, dessicator, oven, evaporating dish.

Reagent: Barium acetate, sodium carbonate, water

Procedure

Weigh approximately 1.5 g (± 0.01 g) of unknown sample into a previously weighed 250 mL

beaker. Add 100 mL of water and stir for about two minutes dissolves. Put the appropriate pre-

weighed filter paper in a Buchner funnel, insert the funnel into a rubber bung and then insert the

funnel with the rubber bung into a separation flask. Attach rubber tubing to the flask and to the

suction part of the tap. Allow the tap to run smoothly to create suction. Wet the filter paper with

water, stir the beaker and gently transfer all the contents of the beaker into the funnel. Rinse the

beaker with water to make sure all the solids are transferred to the funnel. If the filtrate is cloudy,

pour it back into the beaker and re-filter using the same filter paper. Repeat the procedure until a

clear filtrate is obtained. Weigh a dry evaporating dish, Transfer the filter paper with the solid to

the evaporating dish. Dry the filter with the solid in an over for about 15-20 minutes. Place the


evaporating dish in a dessicator and allow to cool. Weigh the cooled solid and calculate the

amounts of reactants.

Divide the filtrate equally into two measuring cylinder. Add several drops of 1M barium acetate

to one cylinder to test for excess sodium carbonate. If a precipitate is formed, then sodium

carbonate is the excess reactant and barium acetate is the limiting reactant. Test the second

cylinder with 1 M sodium carbonate to determine which is the limiting reactant and the excess

reactant. These two tests will help to determine the limiting and excess reactants in the unknown.

Add 5 ml of the solution that produce the precipitate in the test above to the portion of the filtrate

that produced the precipitate. Use another paper to filter the precipitate. Transfer the precipitate

into the oven and dry for 15-20 minutes. Cool the precipitate in a dessicator for and weigh the

precipitate. Record all your results in Table 1.

ATTENTION: All chemical wastes may be flushed down the drain. Use tongs at all times to

hold the evaporating dish.

Questions

1. Does Line 23 add to the same amount of the starting material of the unknown? Why or why not?

2. A 16 g of hydrogen is reacted with 16 g of Oxygen. What is the limiting reactant and how

many grams of water is produced? 2H2 + O2 →2 H2O

3. When solutions of barium chloride and sodium chromate are mixed, a yellow precipitate of

Barium chromate is formed. Describe in general terms how you would determine

experimentally he identity of the limiting reactant?



Table 6

# Test Result

1 Mass of beaker

2 Mass of beaker + unknown

3 Mass of unknown

4 Mass of filter paper #1

5 Mass of evaporating dish #1

6 Mass of dish, paper, BaCO3

7 Mass of Ba CO3

8 Moles of BaCO3

9 Does addition of Ba(C2H3O2) 2to filtrate result in a precipitate?

10 Does addition of NaCO3to filtrate result in a precipitate?

11 What is the limiting reactant?

12 Mass of filter paper #2

13 Mass of evaporating dish #2

14 Mass of dish, paper and precipitate 2

15 Mass of precipitate 2

16 Moles of precipitate 2

17 Total moles of BaCO3 (line 8 + (2Xline 16)

18 Total mole of limiting reactant

19 Total moles of excess reactant

20 Ration of excess reactant to limiting reactant

21 Total mass of limiting reactant

22 Total mass of excess reactant

23 Total mass of all reactants

24

Experiment 7: Complex Ion Formation

Theory

Addition of potassium cyanide (KCN) solution to a solution of silver nitrate (AgNO3) will

produce a white precipitate of silver cyanide (AgCN). The solubility product of silver cyanide is

exceeded for this reaction to occur

KCN + AgNO3 AgCN + KNO3 (1)

The overall reaction can be written as

Ag+ + CN- AgCN

and the solubility product can also be written as

KSP = [Ag+][CN-] (2)

If extra amount of the KCN or other anions is added, the precipitate will dissolve to form a

silver ion complex or complexes of other ions ([Ag(CN)2 or [Ag(X)n where x = other anions).

AgCNs + CN-aq ↔ [Ag(CN)2] aq (3)

On can conclude that a complex ion can be formed when an ion unite with other ions of opposite

charge or neutral molecules (NH3)

Ag+ + 2NH3 ↔ [Ag(NH3)2]aq (4)

If a sulfide solution is added, the complex ion can dissociate into silver ions which is then

precipitated as silver sulfide

[Ag(CN)2-] aq + S2- ↔ Ag+

s + 2CN-aq (5)

The dissociation constant for complex ion (eqn. 5) is

Kdiss = [Ag+][CN-]2/[Ag(CN)2-] (6)

Kdiss for this reaction is 1.0 x 10-21 mol2/dm6 at room temperature

The value of the Kdiss is very small and it shows that the [Ag+] is very small ([CN-] is very large)

that the solubility product of AgCN is not met.

The inverse of Kdiss is called stability or formation constant and is expressed as:

K= 1/ Kdiss = /[Ag(CN)2-]/[Ag+][CN-]2 = 1021 dm6/mol2

The stability of complex ions varies widely among different complexing agent and is

quantitatively expressed as the stability constant.

[Ag(NH3)2]aq ↔ Ag+ + 2NH3

Kdiss = [Ag+][NH3]2/ [Ag(NH3)2]aq = 6.8 x 10-2 mol2/dm6 and

K= 1/ Kdiss = 1.5 x 107 dm6/mol2

The more stable a complex the larger the value of its stability constant

Apparatus: 4x 100 mL beakers, 1x 10 mL measuring cylinder, 3 x 50 mL bakers, stirrer,

stopper, 4 test tubes

Reagent: 0.01M KCN, 0.2 M AgNO3, O.01M sulfide salt,

Procedure

Obtain 15ml of 0.01M KCN and 20 mL 0.2M AgNO3 and 20mL of 0.01M sulfide salt into 3

x100 mL beakers. Prepare the solutions as shown in Table 7.1 into 3 test tubes. Shake each

solution at intervals of 5 minutes for fifteen minutes and leave to stand for precipitation to occur.

This will ensure that equilibrium is reached.

Table 7.1

Solution # Volume of 0.01M KCN

(mL)

Volume of 0.2 M

AgNO3 (mL)

Concentration

(mmoles)

1 5.0 3.0

2 5.0 4.0

3 5.0 5.0

Determine the concentrations of solutions 1, 2, and 3 and record it Table 7.1

[CN-] = (volume of 0.01 M KCN)(0.02M AgNO3)/Total volume of solution

Preparing the Calibration Curve

Decant the supernatant solution in solutions 1-3. Centrifuge the decanted solution for 2 minutes


.

ATTENTION: All chemical wastes may be flushed down the drain. Use tongs at all times to

hold the evaporating dish.

Further Reading

LABORATORY MANUAL

FOR

ORGANIC CHEMISTRY

FIRST EDITION

DEPARTMENT OF CHEMISTRY

BABCOCK UNIVERSITY. ILISAN. REMO. OGUN STATE

2014

Further Readings

1. Quantitative Skoog and West. 3rd Edition. USA

2. General Chemistry Manual. Delaware University Newark Delaware. USA

3. Vogel's Textbook of Practical Inorganic Chemistry (5th Edition) 5th Edition

Appendix A

Table A Table of Solubilities of Some Compounds

Anion Soluble Except Insoluble Except

Acetate, C2O4- All are soluble All are soluble

Arsenate, AsO43- NH4, K, Na

Bromide, Br- Pb, Hg, Ag

Carbonate, CO32- NH4, K, Na

Chlorate, ClO4- All are soluble

Chloride, Cl- Pb (cold water), Hg, Ag

Chromate,

CrO42- NH4, K, Na

Fluoride, F-

Al, Ba, Ca, Fe(III), Pb, Mg,

Sr

Hydroxide, OH- NH4, K, Na

Iodide, I- Pb, Hg, Ag

Nitrate, NO3- All are soluble

Nitrite, NO2- NH4, K, Na

Oxalate, C2O42- NH4, K, Na

Oxide, O2- Ba, Ca, K, Na, Sr

Pechlorate,ClO4- K

Phosphate, PO43- NH4, K, Na

Silicate, SiO4- NH4, K, Na

Sulfate, SO42- Ba, Ca, Pb, Hg, Ag, Sr

Sulfide, S2-

NH4, K, Na, Ba,

Ca

Sulfite, SO32- NH4, K, Na

Thiocyanate,

SCN- Hg, Ag

laboratory manual for inorganic chemistry …€¦ · · 2017-11-23the experiment. no exceptions...

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